Ionisation Energy -> The enthalpy change that occurs when 1 mole of electrons is lost from 1 mole of
gaseous atoms, forming 1 mole of gaseous 1+ ions (or 1+ -> 2+)
Electron Affinity -> The enthalpy change that occurs when 1 mole of electrons is gained by 1 mole of
gaseous atoms, forming 1 mole of gaseous 1- ions (or 1- -> 2-)
Lattice formation Enthalpy -> The enthalpy change that occurs when 1 mole of a solid ionic
compound is formed from its constituent gaseous ions
Lattice dissociation Enthalpy -> The enthalpy change that occurs when 1 mole of a solid ionic
compound is dissociates into its constituent gaseous ions (same value but opposite sign to Latt Form)
Enthalpy of solution -> The enthalpy change that occurs when 1 mole of a solute (solid ionic
compound) completely dissolves in an infinitely dilute solution
Enthalpy of Hydration -> The enthalpy change that occurs when water molecules surround 1 mole of
gaseous ions
Enthalpy of atomisation -> The enthalpy change that occurs when 1 mole of gaseous atoms is formed
from an element in its standard state
Enthalpy of fusion -> The enthalpy change that occurs when 1 mole of liquid substance is converted
into a solid
Enthalpy of vaporisation -> The enthalpy change that occurs when 1 mole of gaseous substance is
formed
These are all under standard conditions!
Born-Haber Cycles:
This is another application of Hess’ law, where we can use an alternative pathway to get the enthalpy of
formation of a solid ionic compound.
This is an example of a Born-Haber cycle for NaCl. There are some steps
you should follow;
Start with elements in standard state at 0 energy
Atomise & then ionise metal atoms (I.E.)
Atomise & then ionise non-metal atoms (E.A.)
Form the lattice with lattice formation enthalpy
The enthalpy change from elements in standard state to the solid
ionic compound is the enthalpy of formation
Once you have drawn out the cycle, including labels for all
enthalpies and state symbols and electrons where necessary, use
the following equation to calculate the enthalpy changes;
ΔfH = Σ All other Enthalpies
Note, as atomisation is the formation of 1 mole of gaseous ions, non-metal atoms will often have to be
multiplied by a number. For example, if you need 2Cl⁻ ions, you would have to multiply the atomisation
enthalpy, and the electron affinities by 2 (All enthalpies involving that ion are multiplied except for the
lattice formation). Watch out for this!
Note, arrows upwards at endothermic and arrows downwards are exothermic (1st E.A. is exothermic, but
2
nd E.A. is endothermic as it requires energy to overcome repulsion between similarly charged species)
Lattice Enthalpies:
The lattice enthalpy of an ionic compound demonstrates how strong the ionic bonding is in that molecule.
The higher the lattice enthalpy, the more energy that is required to separate the oppositely charged ions
from the giant ionic lattice. The magnitude of the lattice enthalpy is affected by 2 factors;
Charge on ions -> larger charges = higher lattice enthalpy
Size of ions -> smaller ions = higher lattice enthalpy (more effective packing)
When comparing molecules to deduce which has the higher lattice enthalpy, look for things both
molecules have in common, then find the feature that isn’t in common and finally state which has higher
lattice enthalpy.
MgCl2 and CaCl2 -> both have Cl-
ions, and both Mg and Ca have 2+
charges, but Mg2+ is a smaller ion that
Ca2+, so MgCl2 has the larger lattice enthalpy.
Bonding Spectrum:
Pure Ionic;
Perfectly spherical ions. However this doesn’t actually exist, as it assumes that the ions are point charges,
there are no forces between the ions, and there are perfectly elastic collisions.
Polarised Ionic;
Ions are polarised (or distorted). All ionic compounds fit into this category, due to the distorted ions in the
ionic lattice
Polar Covalent;
Unequal electron sharing in the covalent bond due to unequal electronegativities of atoms in the bond,
one side of the covalent bond is electron rich (δ⁻), whilst the other side is electron deficient (δ⁺), giving a
polar covalent bond
Non-polar Covalent;
Atoms has the same electronegativity, so the electrons are completely evenly shared in the bond, and
neither side is electron deficient or rich.
Theoretical lattice enthalpies are calculated with Born-Haber cycles, whereas experimental lattice
enthalpies are calculated via experiment. These will never be the same value as the Born-Haber cycle
assumes that ions are perfectly spherical, but the closer the 2 are to one another, the more perfect the
ionic structure of the salt
Enthalpy of Solution:
The term infinitely dilute solution refers to a solution where the addition of more water wouldn’t result in
a further temperature change
ΔsolH = Δlatt dissH + ΔhydH
(Both positive and negative ions)
The enthalpy of solution is affected by ionic radius and ionic
charge. This is because it involves a lattice enthalpy, so is affected
in the same way. Ensure you’re using Latt Diss and not Latt Form!
Entropy:
Entropy is a measure of the disorder of a system. According to the 2nd law of thermodynamics, the
entropy must always increase, so the disorder must also increase.
Solids are the most ordered state, so have the lowest entropy. Liquids
are slightly more disordered, so have a higher entropy and gasses are
significantly more disordered, so have the highest entropy. In
addition, the more moles of a substance there are, the more disorder,
so the higher the entropy.
At absolute 0, the entropy is 0 because no particles have any energy. As the temperature increases, so too
does the entropy. This is because the particles have more energy, so vibrate more, so are more disordered.
There is large jumps in entropy from solid -> liquid and liquid -> gas. The jump from liquid -> Gas > jump
from solid -> liquid. This is because gasses are more disordered than liquids, so have a higher entropy and
so the entropy change is greater.
Entropy Change (denoted as ΔS) can be deduced simply by looking at a reaction. For example, if the states
are more disordered and there are more moles then you can deduce that the entropy must increase, so
the entropy change is positive. However, occasionally, the number of moles and the states are exactly the
same, so you cannot deduce the enthalpy change. Here you must state the enthalpy change is small and
close to 0.
We can also calculate a numerical value for the entropy change using the formula;
ΔS = ΣS (products) – ΣS (reactants)
You add up the entropy of product and reactants and use the above formula. Note, entropy is measured in
J/K/mol, so it is much smaller than enthalpy.
Gibbs Free Energy:
This is used to link the enthalpy change and entropy change of a reaction to determine if the reaction is
feasible or not.
A reaction is only thermodynamically feasible when ΔG ≤ 0. As ΔG is dependent upon temperature, you G is dependent upon temperature, you
must change the temperature to make a reaction which wasn’t feasible a spontaneous (feasible) reaction.
To find the temperature at which a reaction becomes feasible, set ΔG is dependent upon temperature, you G equal to 0 and rearrange to find T.
You have to convert entropy into kJ/K/mol to use in the gibbs free energy equation!
Even if a reaction is thermodynamically feasible, it may not have the Ea to actually occur spontaneously.
You can rearrange the equation into the form ‘ΔG = -ΔST + ΔH’ and plot a graph of ΔG is dependent upon temperature, you G against T. The
gradient is –ΔS (reverse sign to find ΔG is dependent upon temperature, you H) and the y axis intercept is ΔG is dependent upon temperature, you H. The feasibility temperature is
shown by the x axis intercept.